Synthesis of Salicylic Acid and Potentiometric Determination of Its Purity and Dissociation Constant

Synthesis of Salicylic Acid and Potentiometric Determination of its Purity and disassociation Constant - Abstract The purpose of the study is to synthesize salicylic cutting from the ester, methyl radical group salicylate, and determine the panellings disassociation invari fit and probity. The ester was converted to salicylic window glass by immoral hydrolysis. The results were refluxed and recrystallized, to ensure maximum probity, and separate outed, dried, and weighed. The melting occlusion of the harvest-tide was determined using a Fischer-Johns melting rouse apparatus.The deadly indeed dissolved in separate beakers with 95% ethanol and water supply and titrated with 0. 050 M NaOH, previously standardized with kB enthalpy phthalate, through with(predicate) potentiometric titration. The pH after(prenominal) asset of base was mensurable and plotted once morest the vividness of titrant added using trio different plots. Results show a 61. 0% go and the meltin g heyday differed from the conjectural by a range of 3. 11-6. 83%. The pKa calculated was 2. 865, differing from the belles-lettres value of 2. 98, by 3. 86%. The theoretical purity of the sample was deoxycytidine monophosphate. 0%, which differed with the try outal wizard by 1. % the experimental purity is 101. 7%. Potentiometric titration proves to be adequate in the determination of the savage dissociation unremitting and purity of a sample. aside from that, the synthesis proved adequate given the risque purity of the product. - Keywords tart dissociation constant purity melting time period ester INTRODUCTION Potentiometric methods of digest atomic number 18 based on measuring the electric potentiality of electrochemical cells without drawing much, appreciable current. For centuries, potentiometry has been utilize to locate the end run in to the highest degree titration set-ups. Skoog, et al. , 2004). Potentiometric methods offer a myriad of advantages, its main advantage universe its deplorable operational costs. Voltmeters and electrodes atomic number 18, generally, far cheaper than most modern scientific instruments. Models suitable for direct potentiometry in field work, away from the laboratory, atomic number 18 inexpensive, compact, and easy to use. basically nondestructive of the sample, i. e. , insertion of the electrode does non drastically flip-flop the composition of the test base (except for the slight leakage of electrolytes from the reference electrode), electrodes are tellingly free from interferences. stable potential readings are attained fairly rapidly and voltages are intimately recorded as plumps of time. Finally, the wide range of analyte activities, over which some of the lendable indicator electrodes exhibit stable, aroundly Nernstian responses, re delivers an important advantage (Day and Underwood, 1991). Potentiometric titrations involve measuring of the potential of a suitable indicator electrode a s a function of titrant volume. This indicator electrode is speecific to the hydronium ion, H3O+.It offers adavantages over direct potentiometry in that it is not dependent on measuring absolute set of Ecell. This is due to having the measurement based on the titrant volume that pull ins a rapid ex trade in potential near the equation charge. This makes the process relatively free from the juction potential uncertainties as this potential remains approximately constant during the titration process. Instead, the titration results depend heavily on having a titrant with accurately known meanness. The instrument merely signals the end engineer and behaves like a regular chemical indicator.Aside from that, the reference electrode potential need not be known. Most importantly, ionic strength effects are not important in the titration procedure be realize the result is analyte concentration, even off if the electrode responds to activities. The dissociation of a weak monoprotic v irulent is given by the equation HA H+ + A (1) H+A HA where HA is the monoprotic stinging, and A is its conjugate base. The corresponding equilibrium constant for this social disease dissociation is defined as Ka = (2)KaHA A which, when seeking to find H+, rearranges to H+ = (3) front to the equating occlusion, before any titrant was added and the analyte is the only species constitute in ups zealous, the pH of the system is calculated from the concentration of that analyte and its dissociation constant. In the case where a weak monoprotic dot is being titrated with a strong base, subsequent addition of the titrant will cause a response to occur surrounded by the vinegarish and the base. HA + OH H2O + A (4)The presence of the acid and its conjugate base in ascendent will cause the formation of a polisher solution, which are solutions that resist a drastic change in pH, should a strong acid or base be added to the system. At these points, the ramp of a constructed tit ration curve is at its minimum. This is the pre- par point. There is a point during the pre-equality point region wherein both the acid and its conjugate base are present in equal amounts. This occurs when half of the acid has been neutralized, or when the titration is at the half-equivalence point.At this point, the buffering capacity of the system is at its maximum. Aside from that, simplifying Eq. 3 at this point by inspection, the H+ is equal to that of the Ka. Taking the negative of the logarithm of both sides, one will be able to get the relation pH = pKa (Skoog, et al. , 2004). Thus, the dissociation constant whitethorn be computed through determining the pH at half-equivalence point. This experiment will use salicylic acid as the analyte. Salicylic acid, Fig. 1. Salicylic acid is a weak monoprotic acid, capable of releasing the acidic heat content connected to the COOH group.The hydrogen of the phenol group is hard to release because the oxygen is stabilized by resonance. S alicylic acid is a naturally occuring substance, usually produced by plants. It is found in general in the willows leaves and bulwarkk. The pure acid possesses several useful medicinal properties. It is an febrifuge (fever reducer), moderating (pain reliever) and anti-inflammatory (swelling reducer). However, pure salicylic acid makes for an extremely unpleasant medicine. Salicylic acid contains 2 acidic functional groups, the phenolic (C6H5OH) and the carboxylic acid (RCOOH) groups.These groups cause the acid to be an irritating substance that burns the delicate cladding of the mouth, throat, and stomach, hence its esterification to acetylsalicylic acid or aspirin, before ingestion as analgesic (Reed College, 2009). Esterificaton is the process by which a carboxylic acid is trans make to an ester. Esters are organic merges that are derived usually by reacting a carboxylic acid and an intoxi rumpt. The general formula for esters is with the RC=O group derived from the parent carboxylic acid, and the OR group from the parent alcohol.The mechanism for the response of the alcohol and carboxylic acid to form the ester is as follows Fig. 2. Mechanism of Esterification from a Carboxylic acid RCOOH and alcohol ROH Esters whitethorn also be synthesized by reacting the carboxylic acid with other reagents such as SOCl2 to form the acyl chloride, which will thus be treated with an alcohol in pyridine, to esterify it. Esters are among the most widespread of all naturally occuring compounds. Many esters are pleasant-smelling liquids that are credi devilrthy for the fragrant odor of fruits and flowers. For example, methyl butanoate is found n pineapple fossil oil and isopentyl acetate is a constituent of banana oil. The ester linkage is also present in savage fats and other biologically important molecules. The chemical fabrication also uses esters for a variety of purposes. Ethyl acetate, for example, is commonly utilise as a solvent while many another(pren ominal) dialkyl phthalates are used as plasticizers to keep polymers from being brittle (McMurry, 2004). methyl group salicylate is produced by many plants. It was origin isolated from wintergreen leaves, Gaulthea procumbens, and is commonly known as oil of wintergreen. Fig. 3. Methyl salicylateAn ester of salicylic acid and methanol, it masks one of the acidic hydrogens in salicylic acid by substitution it with a methyl (CH3) group. Hence, it is a relatively unreactive compound that does not release salicylic acid efficiently into the body. It is, therefore, not an effective analgesic, or pain-killer. However, it is added to many products, notably for its fragrance, especially root beer and liniments. In order for it to be activated, methyl salicylate must be converted to salicylic acid by organic synthesis, specifically through saponification, a process not unlike that undergone by animal fats to be gravel soaps (McMurry, 2004).In this experiment, salicylic acid will be synthes ized from methyl salicylate by base hydrolysis. Its dissociation constant will also be measured through potentiometric titration. Likewise, the percentage of purity the salicylic acid used in the response will be likewise determined. METHODOLOGY In synthesizing salicylic acid, 1. 2 g of atomic number 11 hydroxide (NaOH) were dissolved in 7 mL water in a round bottom flask.. Half of a mile (0. 5 mL) of methyl salicylate was added to this mixture. The mixture was then efluxed for 15 minutes and cooled to direction temperature. One-milliliter increments of 3 M sulfuric acid (H2SO4) were added until the formation of a clear precipitate, salicylic acid. Half of a milliliter (0. 5 mL) of the acid was added to ensure fill out precipitation of the product. The mixture was then cooled in an ice water privy with a temperature of at most 5C for the reaction to subside. The product was then filtrateed and rinsed with cold water, and recrystallized in water. The solids were then filtered on a pre-weighed filter report and air-dried in the locker.When dried, the solids, along with the filter paper, were weighed and the melting point determined. Two hundred and fifty milliliters (250 mL) of 0. 50 M stock sodium hydroxide (NaOH) solution was then prepared. From this solution, 250 mL of 0. 050 M standard solution was prepared for the titration. This solution was then standardized with 0. 2000 g 99. 95% potassium hydrogen phthalate (KHP) and phenolphthalein as indicator to the pale pink endpoint. Two more(prenominal) trials were done. Three samples of pure salicylic acid were weighed at approximately 0. 000 g and fit(p) in three separate 250-mL beakers. A minimum amount of 95% ethanol was then added to dissolve the acid, and was diluted with 50. 00 mL distilled water. The electrode was scratch line gradational such that it would relate the developed potential to the pH. At this point, the potential would be measured as pH when increments of the titrant are successiv ely added. The base burette, pH meter, and magnetic stirrer were set up according to Fig. 6. Fig. 6. Potentiometric titration set-up A spin bar was place inside the beaker with the sample solution.The electrode was also positioned such that it would not get hit with the spin bar. The pH of the set-up was prototypical measured previous to addition of base. For the first trial, 1. 00 mL of the titrant was added to the analyte and the pH recorded. This was done until the pH was 11. 50. From there, the equivalence point was approximated based on the volume of the titrant that caused a large change in pH. For the two succeeding trials, 1. 00 mL of the titrant was added to the analyte. At 5. 00 mL of the equivalence point, the volume of titrant added was reduced to 0. 0 mL increments. At 3. 00 mL and 2. 00 mL of the equivalence point, it was further reduced to 0. 20 mL and 0. 10 mL, respectively. The titration contnued beyond 5. 00 mL of the equivalence point using 0. 50 mL of titrant u ntil the pH registered was 11. 00. RESULTS AND DISCUSSION Methyl salicylate reacts with a strong base in the following manner Fig. 4. Base hydrolysis of methyl salicylate The methoxy group of the ester is substituted by the hydroxide ion through nucleophilic acyl substitution. The hydroxide attaches itself to the carboxylic carbon.A fast proton transfer from the hydroxide to the methoxy group occurs such that the methoxy gets protonated and leaves the substrate. The sodium ions stabilize the negative charge of the salicylate ion preponderantly found near the two oxygens of the ion. The solution was then refluxed to prevent dismission of material and to prevent the inclusion of impurities in the product. afterward cooling system to style temperature, 1. 00 mL of 3M sulfuric acid (H2SO4) until salicylic acid starts precipitating as a white solid. To complete the precipitation, 0. 50 mL of the strong acid was added to the mixture.At this point, the salicylate ion is protonated and the final product, salicylic acid, forms through the reaction Fig. 5. Protonation of the phenolate and carboxylate groups of the salicylate ion The flask containing the precipitate was then doused in cold water to stop the reaction. Cold water was used in rinsing the solid after filtration to wash out impurities that were water-in dissolvable in the solvent. The solid was recrystallized in hot water. Dissolving the solids in hot water generally increases the solubility of the substances, hence the solids dissolve along with the soluble impurities.The mixture was then allowed to cool slowly. As the solution cools, the solubility of the compound (and of the soluble impurities) passs, the solution becomes saturated with the coveted compound, and the compound begins to crystallize. Because formation of crystals is a highly selective process that usually excludes foreign molecules, only crystals of the coveted compound form. Because the soluble impurities are present in baseer amoun ts, the solution never becomes saturated with the impurities, so the impurities remain in solution even after the solution has cooled.Removing the solution from the crystals thus removes the solvent and the soluble impurities from the desired crystals. A final rinse of minimum ice water ensures the faulting of the impurities still clinging on the surface of the product (Recrystallization, accessed 6 Sept 2010). After it was recystallized, the solid was filtered again in a pre-weighed filter paper. The filter paper used in filtering the solid weighed 1. 0349 0. 0002 g. The mass of the dried product and the filter paper was 1. 3610 0. 0002 g this means that 0. 3261 0. 0003 g was synthesized.Theoretically, with methyl salicylate as the trammel reagent, the mass of salicylic acid should be 0. 5400 g. Computing for the %yield, would give 61. 0%. In the determination of the melting point, the observed range of the melting point was 150-156C. analyze with the theoretical value, 161C, the relative error lies within the range 3. 11-6. 83%. Therefore, the synthesized salicylic acid has a high purity as exhibited by the magnitude of the relative error. This could justify the %yield given that recrystallization might have caused a sledding of material due to the increase of the number of steps involved.It should be historied that in organic synthesis, steps leading to the target molecule do not give 100% yield, hence, increasing the number of steps would decrease the %yield. Aside from that, the low yield could be attributed to the accidental rapid cooling that occurred during the recrystallization process. In a titration curve, there are three important regions pre-equivalence point, equivalence point, and post-equivalence point. For the titration of salicylic acid with NaOH, the pre-equivalence point is characterized by the dependence of pH on the concentration of salicylic acid.Let salicylic acid be HA prior to addition of the base, the pH can be calculated by the concentration of HA and its acid dissociation constant, Ka. When the base is added, HA reacts with the hydroxide ion (OH ) according to the Eq. 4 declared above. Therefore, one has a buffer solution comprised of salicylic acid and salicylate. At equivalence point, all of the acid has been converted into its conjugate base. This base will hydrolyse with water through the reaction A + H2O HA + OH (6) reforming a small portion of the acid. At the equivalence point, the slope of the titration curve is at its steepest.The pH at this point is dependent upon the conjugate base formed and its base dissociation constant, Kb. Beyond the equivalence point, the post-equivalence point, the pH is dependent on the concentration of the excess titrant added. This is because the contribution of the conjugate base is very small and, therefore, negligible compared to that of sodium hydroxide. However, salicylate is still the dominant species of salicylic acid present in solution. A distinction m ingled with equivalence point and endpoint should be established when titrating.Skoog, (2004), states that the equivalence point is the point in titration where the amount of titrant added is chemically equivalent to that of the analyte in solution. The endpoint is a point during titration wherein an observable change signals that the amount of titrant added is chemically equivalent with that of the analyte. The endpoint may not of necessity coincide with the equivalence point, especially in neutralization titrations. It may come earlier or later than the equivalence point, depending on the indicator used. This going pf volume at equivalent and endpoint is known as titration error.In potentiometric titration, no chemical indicator is unavoidable. Instead, the endpoint is characterized by the drastic change in pX, measured by the electrode. X may be H3O+, OH, a cation or anion, or any substance participating in the titrimetric reaction (Day and Underwood, 1991). In this case, the volume of titrant that contributes to the greatest change in pH is the volume required to completely titrate the salicylic acid in solution. Thus, the increments of addition of titrant are narrowed down as the equivalence point is reached because that way, the equivalence point will not be missed.Aside from that, the exact volume of titrant required for complete reaction of the analyte would be detected. Prior to titration and dilution with water, the salicylic acid was dissolved with a minimal amount of 95% ethanol. Salicylic acid is sparingly soluble in water, a polar solvent that can hydrogen-bond with solutes that can hydrogen-bond with its molecules. This is due to the acid having more than 6 carbons increasing its non-polarity, although it has oxygen that can form hydrogen bonds with water (McMurry, 2004). Adding ethanol reduces the polarity of the solvent, facilitating dissolution of the acid.It was noted that some of the acid reprecipitated upon addition of water. Thus, a mi nimum amount of 95% ethanol was again added to dissolve the acid. In the experiment, the electrode was also not lowered while the bar was spinning as air bubbles would adhere on its surface causation an error in the measurement of pH. Constant stirring is important in potentiometric titrations, as in other types of titrimetric analyses. Stirring is done because it will cause the titrant to react not just with the surface of the analyte where it dropped but with the entirety of the analyte solution.This way, the reaction will go to completion and minimal error will be observed. The volumes used for standardization of NaOH with KHP as primary standard are tabulated in put over 2 in the Appendix. Standardization of sodium hydroxide gave 0. 04633 M NaOH. This value was used in the determination of the percentage of purity (%purity) of the salicylic acid sample. For the determination of the equivalence point volume, three plots were done for each trial. The first plot is the graph betw een pH vs. volume of titrant. The second is the first derivative plot with which ? pH/?V was plotted against V, where ? pH and ? V are the change in pH and the volume added between two consecutive readings, respectively, and V is the average volume between two consecutive readings. The third plot is the second derivative plot between ? 2pH/? V2 and V. The ordinate of the graph, ? 2pH/? V2, was obtained by taking the quotient of the difference between two consecutive readings of ? pH/? V, and ? V, whereas V is the average of the two consecutive readings of V. Only two trials were done because of the inadequacy of the time. The graphs for the first trial are a) (b) (c) Fig. 7. Titration Curves for the First exam. (a) pH vs Vtitrant (b) ? pH/? V vs V (c) ? 2pH/? V2 vs V MNaOHVNaOH at eq ptFWsalicylic acid Gsample The graph for the second trial may be found at the Appendix. Either of the three graphs of Figs. 7 and 8 may be used in the determination of %purity for each trial. The equati on used for determining the %purity, specifically for this experiment is %purity = x 100 (7) Thus, the volume of titrant used at equivalence point is required for the calculation.The interpolated set of the volume at equivalence point are tabulated in Table 3 in the Appendix. Calculation of the mean %purity is 101. 7%, having an error of 1. 7%, as the theoretical value is 100. 0%. The pKa is based on the pH at half-equivalence point by virtue of the derivation of Eq. 3. Onle Figs. 7a and 8a may be used for the determination of pKa as the other four do not directly give the pH at each point. The interpolated values for the pH at half-equivalence point is tabulated at Table 3 in the Appendix. The experimental pKa is 2. 865. Therefore, the Ka of the acid, given Ka = logKa (7) or Ka = 10pKa (8) is 1. 3646 x 103. The theoretical pKa is equal to 2. 98. Thus, the pKa value gave rise to a 3. 86% error. Statistical analysis of the results shows the following results Table 1. Statistical ab stract of the Results Range Standard deflexion %purity 15. 7% 11. 1% pKa 0. 21 0. 148 Relative Standard difference (RSD), ppt Confidence Intervals (95% trustingness) %purity 109. 1 101. 7 99. 7 % pKa 51. 8 2. 865 1. 334 The RSD of %purity is relatively large such that the values have low precision.Aside from that, the confidence intervals for the mean is also large such that it almost has the similar order of magnitude as that of the mean. This means that the mean exhibits very low accuracy. With regards to the pKa, the RSD showed a low value, only 51. 8 ppt, implying high precision amongst the values. On the other hand, the mean value shows low accuracy because of the magnitude of the confidence intervals. These errors could be attributed to the number of trials. Due to inadequate time, the group was only able to do 2 trials, one less than the prescribed number of trials.CONCLUSIONS The experiment aimed at synthesizing salicylic acid from methyl salicylate and determining t he acid dissociation constant Ka of the acid along with its purity. Organic synthesis provided a 61% yield of the acid, a relatively low yield. However, the purity of the acid can be classified as high due to the observed melting point ranges precision with the theoretical one the %differenceis only 3. 11-6. 38%. The results of the potentiometric titration show that the salicylic acid used was 101. 7% pure, a value great than the purity, which is 100%.The Ka, expressed as pKa, obtained was 2. 865, 0. 035 units less than the theoretical value, which is 2. 98. Though the %differences are low, the accuracy of the computed values is questionable given that the confidence intervals for the %purity and pKa are 99. 7 and 1. 334, respectively. These errors can be attributed to the number of trials done, which is two, due to the slow stabilization of the pH meter readings that resulted in a long period of titration. Thus, it can be concluded that potentiometric titration is an effective way of determining the acid dissociation constant of a sample.It is recommended that a better pH meter be used in the measurement of the pH and that solutions used be titrated immediately. REFERENCES 1 Day, R. A. and A. L. Underwood. (1991). denary Analysis. Sixth Edition. Prentice Hall, Inc. , USA. 2 Experiment 3 Synthesis of Salicylic Acid from anoint of Wintergreen (Methyl Salicylate). Laboratory Reference Manual Chemistry 201/202. Reed College. Accessed 6 September 2010 http//academic. reed. edu/chemistry/alan/ 201_202/lab_manual/Expt_salicylic_acid/background. html 3 McMurry, J. (2004). Organic Chemistry. Sixth Edition.Brooks/ gelt Thomson Learning, CA, USA. 4 Recrystallization. University of Massachussetts. Accessed 6 September 2010 http// www. chem. umass. edu/samal/269/ cryst1. pdf 5 Skoog, D. A. , D. M. West, F. J. Holler, S. R. Crouch. (2004). Fundamentals of Analytical Chemistry. Eighth Edition. Brooks/ borecole Thomson Learning, CA, USA. APPENDIX Gfilter paper = 1. 0349 0. 0002 g Gfilter paper and dried product = 1. 3610 0. 0002 g Gdried product = 0. 3261 0. 0003 g %yield = 61. 0% Observed melting point range 150-156C %difference = 3. 11-6. 83% 1 standard used potassium hydrogen phthalate (KHP) purity = 99. 95% Table 2. Standardization of NaOH Trial Mass of KHP, g Vtitrant, mL Molarity, M 1 0. 2030 21. 30 0. 04664 2 0. 2175 23. 00 0. 04628 3 0. 2043 21. 70 0. 04608 Average Molarity, M 0. 04633 Table 3. Sample Analysis Trial Gsalicylic acid, g Vtitrant at eq. pt. , mL % purity pH at half-eq. pt. (pKa) 1 0. 1125 16. 50 93. 86 2. 76 2 0. 1031 17. 65 109. 6 2. 97 3 - - - - Average pKa = 2. 865 Average %purity = 101. 7% bandage of the Second Trial (a) (b) (c) Fig. 8. Titration Curves for the Second Trial. (a) pH vs Vtitrant (b) ? pH/? V vs V (c) ? 2pH/? V2 vs V